Synthesising And Testing The Purity Of Acetylsalicylic Acid Biology Essay
Aspirin, also known as acetylsalicylic acid, dates back to 1897, when it was isolated by Felix Hoffmann, a chemist with a German company Bayer. Aspirin is a salicylate drug, often used as an analgesic to relieve minor aches and pains, as an antipyretic to reduce fever, and as an anti-inflammatory medication (1). Acetylsalicylic acid is used in various conditions such as lower back and neck pain, the flu, common cold, burns, menstrual pain, headache, migraines, osteoarthritis, rheumatoid arthritis, sprains and strains, nerve pain, toothache, muscle pain, bursitis (inflammation of a bursa, a fluid-filled sac located around joints and near the bones), and following surgical and dental procedures (2).
Timeline of Aspirin
Hippocrates was left historical records of pain relief treatments, including the use of powder made from the bark and leaves of the willow tree to help heal headaches, pains and fevers (3).
Reverend Edward Stone experimented by gathering and drying a pound of willow bark and creating a powder which he gave to about fifty persons: it was consistently found to be a ‘powerful astringent and very efficacious in curing agues and intermitting disorders’. He had discovered salicylic acid, the active ingredient in aspirin (4).
Willow’s active chemical constituent, Salicin, was identified in 1829 by the French
Pharmacist H. Leroux (5).
Italian chemist Raffaele Piria, isolated Salicin, the active compound in the bark for the first time (6).
Charles Gerhardt, a French chemist, mixed salicylic acid with sodium and acetyl chloride in 1853, creating acetylsalicylic anhydride. The procedure to make this compound was time-consuming and difficult, causing Gerhardt to abandon his project without marketing it (7).
A German chemist named Felix Hoffmann, who worked for a German company called Bayer, rediscovered Gerhardt's formula. Felix Hoffmann made some of the formula and gave it to his father who was suffering from the pain of arthritis (3).
February 27th 1899
Aspirin was patented (3).
At the beginning of the first world war in 1914 the Allies lost their source of aspirin and so they offered large prizes for anyone who could make aspirin (8). A young Australian chemist, George Nicholas, attempted his own production in a small pharmacy in Melbourne. The product was made by reacting salicylic acid with an acrid smelling liquid, acetic anhydride, while being heated. A pure aspirin was produced that more than met the purity requirements of the British Pharmacopoeia (9).
Bayer’s patent to Aspirin runs out (10)(11) .
Professor Peter Elwood conducts a trial into the effects of Aspirin on preventing heart attacks (12).
English scientist Professor Sir John Vane and colleagues, Sune Bergström and Bengt Samuelsson win the Nobel prize for discovering the role of aspirin in inhibiting prostaglandin production (13).
US researchers report preliminary study suggesting that aspirin may delay the onset of senile dementia (14).
US researchers find that Aspirin may protect against Bowel cancer (14).
Aspirin is now being used, or tested, to treat a number of conditions (1).
Salicin is closely related in chemical make-up to aspirin. When consumed, the acetalic ether bridge is broken down. The two parts of the molecule, glucose and the benzylic alcohol then are metabolized separately. By oxidizing the alcohol function the aromatic part finally is metabolized to salicylic acid.(15)
The Benzylic alcohol is circled in orange on the diagram. The Acetalic ether bridge is circled in red on the diagram.
In 1838, Raffaele Piria [an Italian chemist] then working at the Sorbonne in Paris, split salicin into a sugar and an aromatic component (salicylaldehyde) and converted the latter, by hydrolysis and oxidation, to an acid of crystallised colourless needles, which he named salicylic acid (3).
The Hydroxyl group is shown in red, as well as the aldehyde group.
Salicylic acid has the formula C6H4(OH)COOH, where the OH group is ortho to the carboxyl group. It is also known as 2-hydroxybenzenecarboxylic acid. It is derived from the metabolism of salicin (16).
The Carboxylic acid group is circled in blue.
Aspirin (acetylsalicylic acid)
The ester group in the final product Aspirin is circled in red.
Aim 1: To Synthesise Aspirin
My method of synthesising Aspirin involves the reaction between Salicylic Acid and Acetic Anhydride. This is viewed as one of the most popular methods of synthesising Aspirin, as the reagents are available for use within a school laboratory, and the conditions required can be set up with the use of basic heating equipment such as a Bunsen burner and different water baths.
The synthesis of aspirin is classified as an esterification reaction. Salicylic acid is treated with acetic anhydride, an acid derivative, causing a chemical reaction that turns salicylic acid's hydroxyl group into an ester group (R-OH → R-OCOCH3). This process yields aspirin and acetic acid, which is considered a by-product of this reaction. Small amounts of Sulfuric acid (and occasionally phosphoric acid) are almost always used as a catalyst (17).
Synthesising Aspirin from Salicylic Acid is the preferred way of producing Aspirin, as synthesising it directly from the Willow bark would prove very impractical and time consuming, as well as not providing me with the means to compare two commercial methods of making Aspirin.
Raffaele Piria, the Italian chemist who successfully isolated Salicylaldehyde from Salicin, did so by the process of hydrolysis and oxidation.
Hydrolysis is defined as the chemical process of decomposition involving the splitting of a bond and the addition of the hydrogen cation and the hydroxide anion of water (18). In simple terms, it is the addition of water to split a compound. Here is an example of a hydrolysis reaction, where an Ester compound (Ethyl Ethanoate), is split by water in the presence of an acid catalyst, into ethanoic acid and ethanol. The reaction involves the splitting of the Ester bond (-COO) in Ethyl Ethanoate, to give the two compounds that the Ester was made from.
Hydrolysis is a key component of the conversion of Salicin to Salicylic acid. It involves the hydrolysis of Salicin to form the intermediate compound Salicyl Alcohol (Saligenin).
From the following diagrams, I have been able to conclude the reaction mechanism for the hydrolysis reaction. A molecule of water is added to the Salicin complex, which causes the splitting of the Ether and the formation of 2 separate compounds. The –OH from the water molecule joins to the Benzene ring, where the Oxygen atom from the Ether group was previously found in Salicin. The splitting of Salicin into two separate compounds is indicated by the line drawn across the Ether group.
The addition of the –OH group from the water results in the formation of Salicyl Alcohol, which is shown below. The –OH group which has been added is circled in red:
The other product formed from the hydrolysis reaction is glucose. The Hydrogen atom which has been added from the water to form the glucose molecule from the splitting of Salicin is circled in blue:
Glucose is the by-product of the reaction, and Salicyl Alcohol moves onto the next stage of the process, which is oxidation.
Oxidation is the loss of electrons from a compound, and is often indicated by the change in charge of an atom. A simple example is the formation of a Magnesium ion from magnesium metal during the reaction between Magnesium metal and oxygen to form Magnesium oxide.
The half equation is as follows:
Mg ----> Mg2+ + 2 e-
We can see that the Mg has changed oxidation state, increasing to a charge of 2+, which shows it has lost two negatively charged electrons (e-).
The next stage of Piria’s reaction involves the oxidation of Salicyl alcohol to form Salicylic Acid.
The Carbon atom circled, accepts an oxygen atom to form a Carboxylic acid group (-COOH), and is formed from a double covalent bond between the Carbon and oxygen atom. Accepting the oxygen atom, the Carbon atom loses electrons to the oxygen atom to form the double covalent bond, and the Carbon atom increases in charge and polarity to a more positive charge. This is because Oxygen is a more electronegative element, and attracts the electrons from the covalent bond more strongly than Carbon, which gives the Carbon atom a partial positive charge.
The colourless, organic crystalline compound produced is Salicylic Acid, and is the final product of this hydrolysis and oxidation react. However when it was used as a medicine, Salicylic acid was found to irritate the stomach lining and cause painful side effects in patients due to its relative acidity. Chemists tried to create a buffer to eliminate the painful irritation of the stomach lining due to the acidic nature of Salicylic Acid. In 1853 Charles Gerhardt, a French chemist, neutralised Salicylic acid by mixing it with sodium and acetyl chloride, creating acetylsalicylic anhydride.
The buffer produced was Sodium Salicylate, which was then reacted with acetyl Chloride, to produce Acetyl Salicylic Acid, the chemical name for Aspirin.
Nucleophilic Substitution Mechanism
The reaction between Salicylic acid and Ethanoic Anhydride/Acetyl Chloride is found as being a Nucleophilic Substitution reaction. A nucleophile is a chemical species that donates an electron pair to an electrophile to form a chemical bond in a reaction. All molecules or ions with a free pair of electrons are able to act as nucleophiles.
The nucleophile in this case is the Salicylic Acid, as it has a lone pair of electrons on the oxygen atom which is located in the hydroxyl group.
Ethanoic Anhydride is often used as the Ethanoylating agent because it is reactive but not too unpleasant or dangerous. A much more reactive Ethanoylating agent is Ethanoyl Chloride but this is toxic and hazardous to use because it is so reactive (1A).
A basic principle which must be explained in order to understand the reaction mechanism is the fact that the oxygen atoms on the Ethanoic Anhydride are much more electronegative than the Carbon atoms. Electronegativity is a measure of the strength of attraction between an atom and an electron pair. It is known as ‘electron pulling power’.
We can therefore use differences in electronegativity to predict how polar atoms in a compound will be, allowing us to work out the reaction mechanism.
According to Pauling’s electronegativity values, oxygen is a much more electronegative element than carbon. (Oxygen has a value of 3.4, compared to Carbon having a value of 2.6).
This gives the oxygen atoms present in Ethanoic Anhydride a partial negative charge, shown as Oᵟ- , while the Carbon atoms have a partial positive charge, which is shown as Cᵟ+.This means that the Oxygen atom attracts electrons more strongly than the Carbon atom.
The Oxygen atom in the Phenol group in Salicylic acid is also ᵟ-. As well as this, the oxygen in the Phenol group has a lone pair of electrons due to the type of bonding it shows.
This lone pair induces attraction between the ᵟ- oxygen atom in Salicylic acid and the ᵟ+ Carbon group on Ethanoic Anhydride. The Salicylic Acid therefore acts as a nucleophile, and the lone pair on the oxygen atom in the Phenol group of Salicylic acid attacks the ᵟ+ carbon atom in Ethanoic Anhydride.
This results in the substitution of a hydrogen atom from the phenol group, as well as the CH3COO- anion from the Ethanoic anhydride. These then bond, to form CH3COOH, which is known as Ethanoic Acid, and is the waste product of the reaction.
The phenol group is therefore able to esterify with ethanoic anhydride to produce aspirin.
The ‘curly arrow’ mechanism for the reaction is shown below:
During the course of my investigation, I will compare the 2 methods used to synthesise Aspirin from Salicylic Acid. Both my methods involve the use of Ethanoic Anhydride, and follow the same reaction mechanism as shown above. However the differences between the methods are in the reagents and conditions used to synthesise Aspirin. Through my tests for purity, and percentage yield, as well as my comparison of each method with commercially produced Aspirin, I will be able to conclude which method is most efficient.
My methods involve the use of two different catalysts.
The definition of a catalyst is a substance which speeds up a chemical reaction by lowering the activation enthalpy, but can be recovered chemically unchanged at the end (1C).
Due to the physical nature of the reagents, both the catalysis involved in the reactions are Homogeneous.
A homogeneous catalyst is one that is in the same physical state or phase as the reactants. In both methods of synthesising Aspirin, this is the case, as shown below:
Salicylic Acid and Ethanoic Anhydride (in solution)
Concentrated Sulfuric Acid (aq)
Salicylic Acid and Ethanoic Anhydride (in solution)
85% Concentrated Phosphoric Acid (aq)
As shown in the table, both reactants and catalyst are in solution, making the process Homogeneous Catalysis.
The catalyst is added to the solution of Salicylic acid and Ethanoic Anhydride in order to lower the activation enthalpy of the reaction, so in effect, a lower temperature is needed for the reaction to progress. This can be visually described by an enthalpy profile, which shows the energy of the products and reactants along the y axis, and the progress of reaction along the x axis.
The Acid Catalysed acylation Reaction mechanism can be worked out based on the fixed mechanism involved in Esterification. Here are the steps:
Step 1: The proton from the acid attacks the carboxyl oxygen which in turn "pushes" the two electrons in one of the bonds "down”, (it delocalizes the electrons and "spreads them out" between the two
Step 2: The delocalized electrons then, in the presence of the alcohol group, rearrange in such a manner as to create a temporary bond between the two reactants, forming an intermediate.
Step 3: The proton from the –OH group attacks the oxygen in the original -OH portion of the acid forming a positively charged oxygen atom. The Electrons holding the water molecule to the intermediate "flip down", releasing this water, leaving the delocalized intermediate in the end of Step 3.
Step 4: In the final step, the proton added in Step 1 leaves and the electrons left behind flip down "closing" the double bond on the oxygen atom and leaving the ester product.
The mechanism for the homogeneous catalysis that takes place in the synthesis of Aspirin is shown below in skeletal formula, broken down into 6 steps:
File:General Scheme for Acid Catalyzed Nucleophilic Acyl Substitution.png
The formation of an intermediate in acid-based catalysis (as shown in step 2), is a characteristic feature of homogeneous catalysis. The enthalpy profile for the synthesis of Aspirin should look something similar to this:
As can be seen from the Enthalpy profile, there are two humps for the catalysed profile.
One is for each step of the reaction. The intermediate compound then breaks down to give the product and reform the catalyst (1D).
Aim 2: Purifying Aspirin using Recrystallisation
The sample of Aspirin which is directly obtained after vacuum filtration is very likely to contain impurities. These impurities will add to the mass of my sample of Aspirin, which will give me inaccurate results for percentage yield, as well as giving me inaccurate results when I test the purity of my Aspirin.
For these reasons, it is very important that I purify my sample of Aspirin immediately after synthesising it to remove impurities.
The different impurities that could be present in my sample of Aspirin cover a range of possibilities. For example, the sample could contain filter paper from the Hirsch funnel or drying. It could also have been contaminated from the use of lab equipment such as pipettes, beakers and funnels. As well as this, not all of the Salicylic Acid may have been converted to Aspirin, and so there may be some Salicylic acid in the Aspirin sample.
Another important factor that may be overlooked is the fact that throughout the process of synthesising Aspirin, acid is involved, often as well as high temperatures. Although this may not seem a concern, Aspirin can be easily hydrolysed back to Salicylic Acid and Acetic Acid. On their own, water and an ester react very slowly, but the process can be speeded up by catalysis (with acid or alkali) (2A). As in both my methods of synthesis, Acid is used (Sulfuric and Phosphoric). It is possible that Salicylic acid and Acetic Acid may be present in my final product as hydrolysis could have taken place (as shown below).
The process of recrystallisation for purifying compounds is ideal because it removes insoluble impurities as well as soluble impurities.
Aim 3: To Identify my products as being Aspirin
It is essential that after purifying my products, I am able to accurately identify both as being Aspirin. If my products are not able to be identified as Aspirin, then it will render the rest of my investigation useless, because I will not be testing the purity of Aspirin. The main method of determining my product is by its melting point, which I will discuss in further detail later in this investigation.
However it is important to analyse the factors that may cause my product not to be Aspirin, and to do this I must look into Hydrolysis in more detail:
I use acid catalysts in order to synthesise my Aspirin in both methods. When Aspirin is heated with acid in solution, it undergoes hydrolysis. The ester bond breaks and the two compounds join with an -H and an -OH group from the water, to form Salicylic Acid and Acetic Acid. There is the possibility that this may have taken place during my investigation due to human error. The mechanism for this acid based hydrolysis is shown below:
Mechanism of the Acid catalysed Hydrolysis of Esters
An acid/base reaction. Since there is a weak nucleophile and a weak electrophile, protonation of the ester carbonyl allows it to become more electrophilic. Here the acid is used, because of the Bronstead-Lowry Theory that an acid is a proton (H+ ion) donor.
The lone pair on the O atom in the water molecule functions as the nucleophile and attacks the electrophilic C in the C=O, with the electrons moving towards the oxonium ion, creating the tetrahedral intermediate.
An acid/base reaction. The oxygen from the water molecule is deprotonated, and loses a H+ ion.
Another acid/base reaction. The -OCH3 group is lost, but first by protonation, and gains an H+ ion.
Another acid/base reaction. De-protonation of the oxonium ion reveals the carbonyl in the carboxylic acid product and regenerates the acid catalyst.
It is important that I identify my product as Aspirin because if Hydrolysis takes place I may be left with traces of Acetic or Salicylic Acid in my product, which will give me inaccurate results for the rest of my investigation.
I will also use the melting point of Aspirin to determine the chemical nature of my products. Aspirin melts at a fixed temperature, and will therefore allow me to identify my product as Aspirin if it melts at the same temperature.
Aim 4 and 5: To Test the Purity of My Aspirin and compare my Aspirin samples and Methods.
There are a number of different methods of testing the purity of my Aspirin.
Testing the purity of my Aspirin shows how successful the recrystallisation method was at removing the impurities from solution.
The methods are as follows:
Thin Layer Chromatography: This technique is used to separate small quantities of organic compounds.
Titration with known samples of Sodium Hydroxide: This technique is referred to as an ‘aspirin assay’ and is used to compare my Aspirin from both methods, as well as with shop-bought Aspirin.
Iron (III) Chloride Test: This test is used to detect the presence of Salicylic acid, and therefore indicate an impure sample.
Reaction with Sodium Bicarbonate: Aspirin will dissolve in Sodium Bicarbonate solution to produce Carbon Dioxide and sodium acetylsalicylate salt.
These 4 methods will allow me to test the purity of my Aspirin products. As well as this, I will be able to fulfil aim 5 of my investigation and compare the 2 Aspirin samples, therefore finding which method produces the purest sample of Aspirin.
Method 1: Synthesise Aspirin
Equipment and Chemicals
150cm3 conical flask
20cm3 measuring cylinders
Concentrated Sulfuric(VI) acid
Glacial Ethanoic acid
Water bath containing crushed ice
Procedure and Chemical Quantities
After zeroing the scales using a weighing boat, I weighed out 10.03g of Salicylic
I then measured out 20cm3 of Ethanoic Anhydride using a 150cm3 conical flask.
I carefully added the 10.00g of Salicylic Acid to the 20cm3 of Ethanoic Anhydride and swirled the solution in the flask. This took place in a fume cupboard.
Using a pipette, I then added 25 drops of Sulfuric (VI) acid to the flask, swirling the flask after every drop added. The acid acts as a catalyst.
I continued to swirl the flask until impure crystals of Aspirin began to form a ‘crystalline mush’.
I then added 20cm3 of glacial ethanoic acid to the conical flask containing the mixture, which diluted the solution.
I then placed the flask in a cold water bath containing crushed ice to allow the crystals to form.
I then proceeded onto vacuum filtration.
Vacuum Filtration equipment
Clean solvent (e.g. Distilled Water)
250cm3 volumetric flask
I connected the conical flask to a vacuum pump via the side arm. The pump creates a partial vacuum so that the filtrate gets pulled through quickly.
I dampened a piece of filter paper and placed it flat on the vacuum funnel, ensuring the whole area was covered.
I switched on the vacuum pump and carefully poured the Aspirin mixture in to be filtered.
I then disconnected the flask from the vacuum pump before turning it off. This avoided ‘suck back’.
I then placed the filter paper containing my impure Aspirin carefully onto a watch glass and put it into an oven for 10 minutes set at a fixed temperature, to allow the crystals to dry.
Diagram of Vacuum Filtration
Synthesising Aspirin: Method 2
1. I used a balance to weigh a 50 mL Erlenmeyer flask. I then placed 10.08g of salicylic acid in
the flask and weigh again. In the fume cupboard, I then transferred 25.0 mL of acetic
Anhydride from a burette into a 100ml flask, and added it to the flask containing the salicylic acid. I also added 5 drops of 85% phosphoric acid (catalyst) to the flask.
2. I clamped the flask in a beaker of tap water supported on a ring stand over a burner flame. I stirred
the mixture until the salicylic acid had dissolved completely. The water was heated until boiling point, and then the flame was shut off. The flask was kept in the hot water bath for 10 more minutes.
3. While the flask was still in the water bath, I added 10 mL of distilled water to the flask to
decompose any excess acetic anhydride.
4. After a minute, I removed the flask from the water bath and added 20 mL of distilled water. This Let the flask cool to room temperature. As the solution cooled, crystals of aspirin began to appear. The solution was cooled by placing the reaction flask in an ice bath.
5. I then weighed a watch glass and filter paper on the centigram balance.
6. I then proceeded to the vacuum filtration step as shown in the previous method.
Variations between two methods
The main differences between the two methods of synthesising Aspirin are in the catalyst involved.
Both catalysts are acid catalysts, as Method 1 involves Sulfuric Acid as the catalyst, and Method 2 involves Phosphoric Acid as the catalysts. These catalysts may appear similar on the surface, as they both have a similar structure as shown below:
Sulfuric Acid Structure (25)http://www.globalwarmingart.com/images/9/93/Sulfuric_Acid_Molecule_Formula.png
Phosphoric Acid structure (26)http://upload.wikimedia.org/wikipedia/commons/thumb/2/29/Phosphoric_acid.svg/220px-Phosphoric_acid.svg.png
Although the structures appear the same, the most important factor in determining the effectiveness of a homogeneous catalyst such as an acid is to analyse its ability to transfer H+ ions in solution.
This is particularly important in the synthesis of Aspirin because H+ ions are added and removed from the intermediates and Phenol/Carboxylic acid groups as shown in my detailed explanation above.
The pH scale was devised at the beginning of the 20th century by a Danish Chemist called Soren Sorensen. He wanted a simple way to indicate how much acid or alkali was present in a solution. The ‘p’ in pH stands for potens, which is Latin for “power”. So pH is measuring the power of Hydrogen ions in a solution, i.e. its concentration (2B).
pH is defined as:
pH= -lg [H+ (aq)]
An acid which has a higher concentration of H+ ions in solution is a stronger acid, and an acid that has a lower concentration of H+ ions in solution is a weaker acid.
Stronger acids are typically used in dehydration and condensation reactions, as they are able to protonate other compounds.
In a dehydration the –OH group becomes charged (R-OH2+) then it can leave (a double bond is formed) with the acid providing the extra H+ for the alcohol. Then the resulting anion (HSO4- or H2PO4-) can do the elimination reaction creating the double bond. If a stronger acid is used, the transfer of H+ ions will be much faster, and the intermediate will be formed much more quickly, and the reaction is more likely to go to completion. For this reason, I will predict that the stronger acid catalyst is more likely to give the better yield and purity of Aspirin.
To come to a conclusion about which acid catalyst is the strongest, and therefore the strongest H+ ion donor, the pH of both of the acids must be found. Through my research, I found that the pH of an acid solution can be calculated if the molarity of the acid is known.
To avoid any inaccurate values in my calculations, I will set the value of concentration at 0.1 mol/dm3 for both Sulfuric and Phosphoric acid. This allows me to calculate a pH value for each acid which is irrespective of the concentration of the acid in solution, leaving the only variable as the chemical properties of each acid in donating H+ ions.
My calculations are as follows:
H2SO4 (aq) 2H+ (aq) + SO42- (aq)
(The dissociation of Hydrogen ions in Sulfuric acid)
From the equation above, I can see that the molar ratio of Hydrogen ions dissociated compared to Sulfuric acid is 2:1. This means that the concentration of Hydrogen ions will be 2x the concentration of Sulfuric acid, and 2 multiplied by 0.1 is equal to 0.2.
This allows me to put these values into the pH equation to calculate the value for pH.
H3PO4 (aq) H+(aq) + H2PO4- (aq)
(The dissociation of Hydrogen ions in Phosphoric Acid)
From the equation above, I can see that the molar ratio of Hydrogen ions dissociated compared to Phosphoric acid is 1:3. This means that the concentration of Hydrogen ions will be 1/3 of the concentration of Phosphoric acid, and 1/3 x 0.1 is equal to 0.03 recurring.
This allows me to put these values into the equation to calculate the value for pH.
From these calculations, I am able to conclude that the pH of Sulfuric acid in 0.1 molar solution is stronger than the pH of Phosphoric acid, and is able to transfer more H+ ions therefore, so theoretically should be the more effective catalyst in assisting Condensation reactions, and therefore should provide a better, purer yield of Aspirin. So Method 1 should theoretically be the most effective method.
Percentage Yield is used to see if a reaction is economically viable.
Using percentage yield will allow me to determine how effective the chosen method is at synthesising Aspirin. To calculate percentage yield for the reaction, I must use the overall equation for the reaction which begins with Salicylic Acid.
C7H6O3 + C4H6O3 C9H8O4 + C2H4O2
Salicylic Acid + Ethanoyl Chloride Aspirin + Acetic Acid
The only compounds I need to use for this calculation are highlighted in bold, and are Salicylic Acid and Aspirin.
The equation for Percentage yield is as follows:
% yield= Actual Mass of Product
Theoretical Maximum mass of Product x 100
The Relative Molecular Mass of Salicylic acid is 138 and the Relative Molecular Mass of Aspirin is 180.
From this information, I can work out the Theoretic maximum yield of Aspirin.
In my investigation, I used 10.03g of Salicylic Acid while producing 11.12g of Aspirin.
The theoretical maximum yield of Aspirin from Salicylic Acid in terms of the quantities I used is calculated:
180/138 x 10.03 = 13.08g of Aspirin
I use my calculated value for Theoretical yield with my value for my experimental (actual) yield to calculate percentage yield for Method 1:
% yield = 11.12/13.08 x 100
% yield = 85.02%
Method 1 has a high yield of Aspirin.
I will use the exact same process for method 2.
Using method 2, I used 10.08g of Salicylic Acid which produced 10.97g of Aspirin.
Theoretical Maximum yield = 180/138 x 10.08 = 13.14g of Aspirin
% yield = 10.97/13.14 x 100 = 83.48%
Method 2 has a slightly lower yield of Aspirin compared to method 1.
Note: There may still be impurities in the Aspirin, which mean that these results cannot be taken as final until the Aspirin has been purified. Impurities such as filter paper and lab contamination, as well as locked-in moisture may cause more mass than the Aspirin alone.
Method 2: Purifying the Aspirin- Recrystallisation
Recrystallisation is a technique used to purify solid crude organic products. The mixture should contain mainly one product, with small amounts of impurities. It works on the principle that only the desired compound will dissolve to an appreciable extent in a suitable hot solvent (2B).
Recrystallisation depends upon the fact that different substances have different solubilities in different solvents. In the simplest case, the unwanted impurities are much more soluble than the desired compound. The impure sample is dissolved in the minimum volume of hot solvent to form a saturated solution, then as the solution cools slowly, crystals of the desired compound form and can be collected by filtration. The soluble impurities remain in cool solution and so pass through the filter paper with the solvent.
Choosing a solvent
The choice of solvent is critical to the purification by recrystallisation, but there is no simple way
to know which solvent will work best. However, some general principles of solubility will assist in this
choice. Crystallisation depends primarily on solubility relationships. In most instances the solubility of a compound in any solvent increases markedly with temperature.
The solubility of crystalline organic compounds depends on the functional groups that are present
and the polarity of the solvent to a very large extent. In this context the expression "like dissolves like" is a very useful principle. Compounds with groups such as -OH, -NH-, -CONH- and others are usually more soluble in solvents such as alcohols or water than in hydrocarbons. Since Aspirin contains an –OH group as well as a –COO ester group, the solvent must be a polar solvent in order for my sample to dissolve.
This method is effective because the aspirin will dissolve in the water due to hydrogen bonding, (which is explained in detailed below), however, solid impurities will not dissolve. When the solution is filtered, the solution will pass through the filter paper but any solid impurities will not, this means that solid impurities will be removed.
One essential characteristic of a useful solvent is that the desired compound must be considerably more soluble in the solvent when it is hot than when it is cold. To determine whether or not
the solvent fulfils this requirement, I did a preliminary test. The steps are as follows:
I placed a 25mg spatula tip of my Aspirin samples into a small separate test tubes, and added 3 drops of water to each.
On observation whilst agitating the flask, I noted that neither of the Aspirin samples dissolved in cold water.
I placed a thermometer into each of my test tubes, and proceeded to heat the tubes using a Bunsen burner whilst monitoring the temperatures of each solution.
I observed that Method 1 Aspirin and Method 2 Aspirin dissolved in the water between temperatures of 70 to 90 degrees Celsius.
I then cooled the solutions by placing them in a crushed ice water bath at 5 degrees Celsius, and observed that as the temperature of the test tube began to drop past 30 degrees my Aspirin began to recrystallise until it had fully recrystallised at 10 degrees Celsius.
After this trial, I was able to conclude that water is an appropriate solvent to use to recrystallise my Aspirin, as the Aspirin dissolved when hot, but remained solid when cold. This means that any insoluble impurities as well as soluble impurities will be filtered out of the solution.
To understand the principle by which Aspirin is able to dissolve in water and then recrystallise, hydrogen bonding must be looked at in more detail:
Hydrogen bonding occurs when there is a large dipole between a Hydrogen atom and either a Fluorine, Oxygen or Nitrogen atom which it is bonded to. The result is that the molecule has atoms which have partial charges. For example, in a water molecule, the Oxygen atom has a partial negative charge (δ-), whereas the Hydrogen atoms have partial positive charges (δ+). This is the same in any other molecule containing these atoms, and the Hydrogen will always have the partial positive charge (δ+).
To understand why the different atoms have partial charges, the principle of electronegativity must be examined. Although I have already explained this principle, it is equally as important to understand it in the dissolving context.
Electronegativity measures the ‘electron pulling power’ of a particular element. Highly electronegative elements attract electrons very strongly compared to weak electronegative elements. For example, in the Hydrogen Chloride molecule, Hydrogen has the lowest electronegativity (shown by the Pauling scale which features earlier in this report), and Chlorine has a higher electronegativity. As Chlorine is more electronegative, it attracts the electron in Hydrogen’s outer shell as well as its own, and so the electron pair moves slightly to the side of Chlorine, giving it a partial negative charge. This is represented by the partial charges shown in the diagram below:
Also shown in the diagram is the attraction between the Chlorine and the Hydrogen atom. This is because the opposite charges attract each other, and form a weak instantaneous dipole-induced dipole bond.
This can be seen in even more electronegative atoms such as Fluorine:
The Fluorine atom attracts the bonding electrons more strongly, giving it a partial positive charge, and the Hydrogen a partial negative charge.
This dipole principle can be seen in the –COOH group on the Aspirin molecule:
δ- δ+ http://upload.wikimedia.org/wikipedia/commons/thumb/6/67/Aspirin-skeletal.svg/336px-Aspirin-skeletal.svg.png
The oxygen atom is more electronegative than the Hydrogen atom and so a dipole is formed on the –COOH group of Aspirin. As well as this, Hydrogen has only 1 outer shell electron compared to 6 in oxygen. Due to this, the oxygen atom only donates 2 of its electron pairs to the bond with the Carbon atom and the bond with the hydrogen atom, which leaves it with a lone pair of electrons. This means it is able to Hydrogen bond with another molecule.
An example of this which relates to recrystallisation is the hydrogen bonding between Aspirin and water as shown below:
The lone pair on the oxygen atom gives it a partial negative charge which attracts the partial positive charge of the hydrogen on the water molecule. This forms a Hydrogen bond, which is indicated on the diagram as a series of lines.
This Hydrogen bonding between the Aspirin and water molecules is able to overcome the solute-solute attraction and so Aspirin is able to dissolve in water.
(Chemical Ideas 13.1)
The method is as follows:
Samples of Aspirin from both methods
A suitable solvent (Hot water or Ethanol)
250cm3 Conical flask
Vacuum filtration equipment- 250cm3 Volumetric Flask and tubing
I proceeded to heat the minimum amount of water needed to dissolve my aspirin, using a Bunsen burner following the correct safety precautions. I then dissolved my sample of aspirin in the hot water. This procedure was done in the 250cm3 beaker.
I then poured the solution into the 250cm3 conical flask through filter paper using a funnel. This filtered out any insoluble impurities which would not dissolve in water.
I then put the conical flask containing my dissolved Aspirin into an ice bath cooled to between 5 and 10 degrees. This allowed the Aspirin to recrystallize, whilst the
I then used vacuum filtration to filter out any soluble impurities, washing once with cold water.
Finally, I scraped off the crystals from the top of the Buchner funnel and left on filter paper, and put in an oven for 15 minutes to dry.
This diagram explained the filtration stage of recrystallisation, when the pure product in an impure solution is poured into the conical flask through filter paper, and the pure product is collected (27).http://www.sciencequiz.net/lcchemistry/02_bonding/benzoic_acid/benzoic.png
Aim 3- Identify my product as Aspirin using Melting Point.
When an organic solid is heated, the heat energy that’s added to the substance is translated into kinetic energy – allowing the molecules to move. The more mobile molecules are able to partially overcome the intermolecular attractive forces which keep them adhered rigidly in place in the highly-ordered structure of the crystalline “lattice.” The individual molecules can move more freely in the liquid state, and the interactions between them become much weaker.
The melting point of a substance is the temperature range over which the first crystal of a solid just starts to melt and the last crystal completes its melting.
A melting point range is very narrow for pure solids (usually just 1 – 2 degrees), and it is a physical property which is unique to the particular compound. Therefore melting point can be used to identify pure compounds in their solid state. The presence of even a small amount of impurity will lower a compound’s melting point by a few degrees and broaden the melting point temperature range. Because the impurity causes defects in the crystalline lattice, it is easier to overcome the intermolecular interactions between the molecules, and therefore, the compound will have a lower melting point with a larger range.
In pure Aspirin, the molecules will all the the same and so will form a regular crystalline arrangement, and therfore more intermolular bonds form, as there are lots of points of contact between the molecules. This gives pure Aspirin a higher melting point due to more intermolecular forces between its molecules. This can be seen from the diagram below:
In impure Aspirin, the impurities will interrupt the crystalline structure of arrangement as they are different sizes and do not contribute to the regular structure of the Aspirin molecules.
The impurities cause a ‘kink’ in the structure, and an irregular arrangement. Due to this irregular arrangement, there will be less intermolecular bonds formed between the molecules, as they do not pack as closely together. This gives impure Aspirin a lower melting point compared to pure Aspirin.
This can be seen from the diagram below:
It is also important to look at melting point in terms of the structures involved:
The structure of Aspirin is important in explaining its melting point:
The Carboxylic acid (-COOH) group on the Aspirin allows molecules of Aspirin to form hydrogen bonds with each other (this is explained in the Hydrogen bonding section). This means that the melting point of aspirin will be relatively high. Permanent dipole – permanent dipole bonds may also be formed from the C=O bonds on the aspirin molecule. These bonds occur when one atom is more electronegative than another and a dipole is set up (see electronegativity). These opposite charges attract and cause a bond to form.
The Hydrogen bonding in Aspirin is shown below:
However, the structure of Salicylic Acid is different. Salicylic Acid contains a –COOH group, but as well as this contains an –OH group. This gives Salicylic Acid more opportunities for Hydrogen bonding compared to Aspirin. This is because bonding can come from both the –COOH group and the –OH group. For this reason, there are more Hydrogen bonds formed between Salicylic acid molecules, and the melting point is higher, because more of these Hydrogen bonds will require more energy in the form of a higher temperature to break.
(Bonding between Salicylic Acid molecules)
Permanent Dipole-Permanent Dipole Bond
Standard Melting Points
Salicylic Acid (2-Hydroxybenzoic acid)
(This information is from the Royal Society of Chemistry at www.rsc.org)
Using this information, I will be able to compare my results for melting point to test the purity of my products.
If my products melt inside the range given, it will indicate that my Aspirin samples are pure.
Bunsen burner and safety equipment
Melting Point equipment (Fisher-Johns equipment)
Both Method 1 and Method 2 samples of Aspirin
Sample of Salicylic Acid
Sample of Impure Aspirin
Sample of commercial Aspirin
Using a Bunsen burner, melt the end of 5 capillary tubes whilst on a semi-blue flame. This step must be done carefully, as if the tube is heated too much it will bend.
Carefully transfer the samples of Salicylic Acid, both samples of Aspirin, impure Aspirin and commercial Aspirin each into individual capillary tubes.
The tubes must be agitated and flicked to ensure the compound has reached the bottom of the tube.
Set up the melting point apparatus by ensuring the thermometer is securely in the machine.
Place one of the capillary tubes into the apparatus and turn on the heat, observing through the magnification the crystals of product.
As the temperature rises, note the temperature range at which the compound begins to melt and when it has turned completely liquid.
Record this temperature and compare with the standard value for melting point.
Repeat this procedure with the other capillary tubes, and compare results with the standard values for melting point.
Iron (III) Chloride Test
Iron (III) Chloride is added to a solution of Aspirin in order to determine if it had hydrolysed back into Salicylic Acid. It is therefore, a purity test.
2g samples of my Aspirin (both methods)
2g sample of Salicylic Acid
2g sample of pure (commercial) Aspirin
Solution of Iron (III) Chloride
I weighed out 2g of my sample of Aspirin using a scale after zeroing using an empty weighing boat, and placed the 2g into a 250cm3 beaker.
I then poured the minimum amount of hot water needed to dissolve the Aspirin, and agitated the beaker until it had fully dissolved in solution.
I then put 3 drops of Iron (III) Chloride solution into the beaker and swirled until the colour change was visible.
I observed and recorded the colour change.
I repeated the procedure with all of the 2g samples shown in the equipment list.
Iron (III) (Fe3+) is a transition metal. A transition metal had distinct properties and is defined as:
Transition Metal: an element which forms at least one ion with an incomplete d sub-shell containing at least one electron.
The general properties of the transition elements are
They are usually high melting point metals.
They have several oxidation states.
They usually form coloured compounds.
They are often paramagnetic.
The Bohr model of the atom proposed by Nicholas Bohr shows that atoms exist in quantised energy levels or shells. These shells are split into sub-shells, which are used to order the period table, and so the periodic table allows us to predict the electron configuration of an atom.
The four different types of orbitals s,p,d, and f have different shapes and one orbital can hold a maximum of two electrons. The p, d, and f orbitals have different sublevels unlike the s orbital and therefore can hold more electrons. Each of the these subshells can hold a different maximum number of electrons.
When the subshells are being filled, they fill in a particular order which can be seen across the periodic table. When written out, the order is as follows:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p.
Each shell has a fixed number of maximum electrons it is able to hold (as mentioned above), and these are:
Maximum number of electrons
This trend can be seen across the periodic table:
As shown, the periodic table is divided into 4 sections. The transition metals are shown in the red section of the table, the d subshell. It is possible to predict the electron configuration of transition metals by looking at the periodic table and counting across. This can be shown using the different ionic states of Iron.
Iron’s electronic configuration is 1s22s22p63s23p63d64s2.
Each number (those in ascending order) represents a shell, each letter represents a sub-shell (seen in the periodic table) and the superscripted numbers represent the number of electrons each sub shell is holding.
By adding all the superscripted numbers together, we get to a total of 26, which shows that Iron as an unionized transition metal has 26 electrons altogether.
Iron is able to exist as ions, and it loses electrons to gain a positive charge which can be shown in a configuration diagram. In this case, we are able to see the arrangement of the individual electrons in each subshell. Iron loses 2 electrons to become Fe2+, and 3 electrons to become Fe3+. This is shown below:
Each arrow represents one electron, with two in each square making a pair. The 3d shell fills up first, going along, with one electron filling up each shell (as seen in Fe3+). Electrons are lost from the 4s shell first, as seen in the ions of Fe. The different oxidation states of Iron are shown in the table, as 2 and 3. The different oxidations of Fe allow it to form complex molecules with different ligands.
A complex consists of a central metal atom or ion surrounded by a number of negatively charged ions or neutral molecules possessing a lone pair of electrons. These surrounding anions or molecules are called ligands. (Chemical Ideas 11.6)
The Iron (III) Chloride test is used to distinguish between aspirin and salicylic acid. Salicylic acid contains a phenol group whereas Aspirin does not. The Iron (III) Chloride reacts with the phenol group to form a ligand complex, and Hydrogen Chloride.
The Ligand substitution reaction is as follows:
FeCl3 + 6C7H6O3 [Fe (C7H5O3-)6]3- + 3H+Cl-
The Phenol group is able to form a dative covalent bond with the central metal ion (Fe), and form a ligand. However, this is not possible in Aspirin because it does not have a Phenol group with a lone pair of electrons on the oxygen atom, and so cannot form a dative covalent bond with Fe.
For this reason, Aspirin does not react with Iron (III) Chloride, however Salicylic Acid does.
This allows me to see if my Aspirin is pure and the –OH group has been completely Acylated.
The complex formed has an octahedral shape as there are 6 areas of electron density around the central Fe3+ ion, as there are 6 dative covalent bonds. This means that there are bond angles of 90°.
90° bond angle.
This is the 3D complex formed when Iron (III) Chloride reacts with Salicylic Acid.
The complex forms as a result of the dative covalent bond formed between the (C7H5O3-) ligand and the central metal ion (Fe3+).
When the ligands bond with the transition metal ion, there is repulsion between the electrons in the ligands (C7H5O3-) and the electrons in the d orbitals of the metal ion (Fe3+).This raises the energy of the d orbitals. However, because of the way the d orbitals are arranged in space, it doesn't raise all their energies by the same amount. Instead, it splits them into two groups.
When 6 ligands are arranged around a transition metal ion, the d orbitals always split into 2 groups in this way - 2 with a higher energy than the other 3.
The size of the energy gap between them (shown by the blue arrows on the diagram) varies with the nature of the transition metal ion, its oxidation state (in Fe3+ this would be +3), and the nature of the ligands.
When white light is passed through a solution of this ion, some of the energy in the light is used to promote an electron from the lower set of orbitals into a space in the upper set. The complimentary colour to that absorbed is seen. Therefore the colour change is caused by a split in energy levels.
In the case of the Salicylic Acid complex, it absorbs in the green-yellow wavelength, and emits purple/violet light. This is the colour seen in the test for a Phenol, and if the solution formed is purple when Iron (III) Chloride is added it means that my sample of Aspirin contains Salicylic Acid impurities due to hydrolysis of Aspirin.
Here is the diagram of the complex formed showing the dative covalent bonds between (C7H5O3-) and Fe3+. Each arrow represents a dative covalent bond formed between the lone pair on the Oxygen of the ligand and the central metal ion.
Thin Layer Chromatography
Chromatography is an important analytical technique used to separate mixtures of substances into their components/compounds.
There are different types of Chromatography, and these depend on the equilibrium set up when a compound distributes itself between two phases. These phases are the stationary phase and the mobile phase. The stationary phase stands still while the mobile phase moves over it. In simple terms, the T.L.C paper acts as the stationary phase while the solvent acts as the mobile phase. This is because the T.L.C paper holds the spots of substance to be separated while the solvent moves up the paper throughout the Chromatography process. The equilibrium is different for different compounds because they all move along the mobile phase at different speeds, and have different affinity for the mobile phase.
Chromatography acts as an essential technique in determining the purity of my Aspirin samples. I will be able to visibly see the differences, and compare using Rf values (explained later). Chromatography will also allow me to identify any impurities in my products, again using Rf values.
To understand the principle behind Chromatography, the technique must be looked at from an equipment perspective:
T.L.C is done using a specific paper which has important properties. Paper is made of cellulose fibres, and cellulose is a polymer of the simple sugar glucose. The key point about cellulose is that the polymer chains have -OH groups protruding from all sides due to the 6 Carbon structure. To that extent, it presents a similar surface as silica gel or alumina in thin layer chromatography. A key feature is that Cellulose fibres attract water vapour from the atmosphere as well as water that was present when the paper was made. The paper is essentially cellulose fibres with a very thin layer of water molecules bound to the surface.
It is the interaction with this water which is the most important effect during paper chromatography.
Given two compounds which differ in polarity, the more polar compound has a stronger interaction with the silica/T.L.C surface and is therefore more capable to remove the mobile phase from the binding places. Therefore, the less polar compound moves higher up the plate (which results in a higher Rf value). If the mobile phase is changed to a more polar solvent or mixture of solvents, it is more capable of dispelling solutes from the silica binding places and all compounds on the TLC plate will move higher up the plate. For this reason, I will use Ethanol as my solvent in the T.L.C experiment, as this will give me a higher solvent front and more accurate Rf values.
The diagram shows the behaviour of compounds A and B as the solvent passes between the particles of Silica gel in the thin layer.
The appearance of T.L.C paper after the procedure is a series of spots- one for each compound in the mixture.
The distance travelled relative to the solvent is a constant for a particular compound as long as you keep everything else constant (e.g. the type of paper and the exact composition of the solvent).
The distance travelled relative to the solvent is called the Rf value (retention value). For each compound it can be worked out using the formula:
For example, if one component of a mixture travelled 9.6 cm from the base line while the solvent had travelled 12.0 cm, then the Rf value for that component is:
Due the fact that the solvent front is always larger from the distance travelled by the solute, Rf values are always between 0 – 1. Rf values do not have units since it is a ratio of distances.
In the diagram shown, Caffeine is used as an example. The Rf for substance A would be B/C in this case.
(Chemical Ideas 7.3)
Pencil and ruler
Samples of both Aspirin products
Sample of Salicylic Acid
Sample of pure (commercial) Aspirin
Cling film/Watch glass
Solvent for Chromatography: Cyclohexane, ethyl Ethanoate, ethanoic acid (200:100:1 ratio)
Bunsen Burner apparatus
To begin, I set up a Bunsen burner apparatus, and heated the centre of a capillary tube briefly in a closed flame. This melted the centre of the capillary tube slightly, and when this was broken in two, gave an even thinner and more precise capillary tube. I repeated this technique, making 4 thin capillary tubes.
I then made up standard solutions of both of my Aspirin samples, pure (commercial) Aspirin and Salicylic Acid, using Ethanol as the solvent.
I drew a faint pencil line about 1cm from the bottom of the T.L.C plate.
I carefully spotted my solutions on the pencil line separately, labelling each to ensure they are easily identified afterwards.
I then placed the T.L.C plate into a beaker with a small amount of the Chromatography Solvent detailed in the equipment list. I ensured that the solvent did not touch or rise above the pencil line.
I covered the beaker with a watch glass and allowed the solvent to rise.
When the solvent had almost reached the top, I took the T.L.C plate out of the beaker and marked the solvent front using a pencil, and left it to dry.
I then used a UV light to observe the circles produced by my compounds, and compared my Aspirin samples to the pure Aspirin and Salicylic Acid samples.
I measured and calculated Rf values, and compared them to the known data values to determine purity.
An Aspirin assay is an analytical test for the purity of Aspirin. It involves the titration of Aspirin with Sodium Hydroxide (NaOH) a base.
As seen from the diagram below, the equipment is set up as a standard titration.
As can be seen from the diagram, the apparatus is set up as shown. The NaOH solution is fed into the Erlenmeyer flask from the burette. The Erlenmeyer flask contains the samples being titrated in their standard solutions
In my first attempt at the Aspirin Assay, I found that my titre was much larger than the given results for pure Aspirin. This was unexpected and questioned the purity of my samples. However, after some research, I was able to find that the problem of the dual acidic nature of Aspirin was causing the anomalies. After amending my method, I was able to correct this error and improve the accuracy of the titration analysis. The error is explained below.
When aspirin is titrated, first the salicylic acid is titrated and then, as the Ester bond breaks, the acetic acid is also titrated, but this breakdown of the aspirin is slow. In theory, both titrations should take equal amounts of base since one mole of aspirin is made from one mole of salicylic acid (neutralized in the first titration), and one mole of acetic acid (neutralized in the second titration). Unfortunately this is not the case. The breakdown of aspirin requires an excess of base to ensure that the reaction is complete. So when aspirin is titrated it is done in steps:
First, enough base is added to titrate the salicylic acid to a phenolphthalein end point.
Next, an equal amount of base is added to titrate the acetic acid that will be released, but since this reaction requires excess base for the reaction to occur, extra base is added.
To overcome this problem, a known excess amount of base is added to the sample solution and a HCl titration is carried out to determine the amount of unreacted base. This is subtracted from the initial amount of base to find the amount of base that actually reacted with the aspirin and hence the quantity of aspirin in the sample.
This procedure allows me to have a much more accurate titre, as any excess base will be removed from the titration results.
The following skeletal equation represents the overall reaction taking place during the titration:
+ NaOH(aq ) + H20(l)
The equation for the reaction is as follows:
C9H8O4(s) +NaOH (aq) ----->C9H7O4(s)-Na++H2O (l)
The –COOH group from the Aspirin loses an H+ ion to form the carboxylate salt group, with a negatively charged oxygen atom on the outside. This negatively charged oxygen atom is attracted to the positively charged sodium ion from the base, and forms the sodium salt. The H+ ion that is lost is attracted to the OH- ion which acts as a nucleophile, and accepts the proton to form H20, which makes the overall reaction a neutralisation reaction. This is because the Aspirin is slightly acidic , and a salt and water are formed.
The “end point” of an acid-base titration is determined by a colour change. In this case, it is the first sign of a permanent pink tinge.
The pink colour is essential in explaining the chemistry behind the assay. The titration is done to determine the amount of Aspirin in the sample being tested. Therefore it is testing the purity of the sample, as the other compounds present will be impurities. Aspirin is classed as a weak acid because of the carboxylic acid group which gives the molecule a relatively stable anion (COO-) which means that it can readily donate a proton (H+ ion). This corresponds with the Bronstead-Lowry theory that “an acid is a proton donor while a base is a proton acceptor.”
Phenolphthalein indicator plays an integral role in determining the end point. In acidic conditions, the phenolphthalein indicator is colourless, along with the standard solution of Aspirin. However in basic conditions, the phenolphthalein indicator is pink. When the Aspirin sample is in solution, the solution is acidic because of the resulting H+ ions given off, and the indication remains colourless. However, when sodium hydroxide is added, it neutralises the acid by the reaction between the H+ ions from Aspirin and the OH- ions from the NaOH to form a neutral solution of water. When the next drop of Sodium Hydroxide is added, theoretically, the solution will turn pink because the conditions will have changed to basic conditions, and the indicator is pink in basic conditions.
However, in practice, as the NaOH is added, the solution shows tinges of pink which disappear as the conical flask is agitated. This is because there are still H+ ions in the Aspirin solution which have not yet been neutralised, and so the solution has not yet been completely saturated with OH- ions enough to completely neutralise it. For this reason, it is important to wait until the first pink tinge remains. This ensures all the H+ ions have been reacted with OH- ions, and the solution is no longer acidic. This means the next drop of Sodium Hydroxide will turn the solution pink, and the titration is stopped, and the titre taken.
The actual colour itself is due to a change in the sub-shell configuration of phenolphthalein resulting in green light being absorbed and the complimentary colour (pink) being emitted (29).
The complimentary acid and base pairs of phenolphthalein can also be used to explain the colour change.
Representing phenolphthalein in acid conditions as Hx, I can construct equations to show the colour changes:
An equilibrium reaction is set up:
Hx(aq) H+(aq) + x-(aq)
When Sodium hydroxide (NaOH) is added, the OH- group reacts with the H+ ion to give water:
Hx(aq) + OH-(aq) x-(aq) + H2O(l)
Since the concentration of H+ has been reduced, the equilibrium shifts towards the side of the products. This causes more x- to be formed and as a result, a stronger pink colour is formed in the solution.
(2C)(Chemical Ideas 8.1)
Mortar and Pestle.
Burette and clamp stand.
6 x 100cm3 conical flasks.
3 x 10cm3 measuring cylinders.
Recrystallised Aspirin samples from both methods.
Pure Aspirin (provided).
0.1 mol dm-3 Sodium Hydroxide solution.
I began by grinding up an Aspirin tablet from the commercial Aspirin provided using a mortar and pestle.
I then zeroed the weighing scales using a weighing boat, and weighed out 1.49g of the pure Aspirin, before placing it into a 100cm3 conical flask.
I then did a similar process and weighed out 1.50g of my own samples of Method1 and Method 2 Aspirin, and placing them into separate 100cm3 conical flasks ready for titration.
I also weighed out 1.48g of the pure Aspirin provided by the lab and transferred this sample into a 100cm conical flask.
Next, I transferred 10cm3 of each solution into different 100cm3 conical flasks, which were to be used for the titration itself.
I then carefully poured the 0.1 mol dm-3 Sodium Hydroxide through a funnel into the burette, ensuring it reached the 0 mark at the top of the burette.
I added 3 drops of Phenolphthalein indicator to the pure Aspirin solution, and swirled the flask to ensure the indicator had spread through the solution.
I then ran the Sodium Hydroxide from the burette into the pure Aspirin solution until the first tinge of pink remained.
After doing this, I filled a burette with 0.1 mol dm-3 Hydrochloric acid until the solution had turned colourless again, and marked the volume added.
This value for HCl acid volume was subtracted from the original titre and the final titre was found.
I then recorded my titre as a rough titre, and repeated the procedure until I had 3 concordant titres (range of ±0.20cm3).
I repeated the titration procedure detailed above with both my samples of Aspirin, and the Aspirin tablet, finding an average titre for each.
Nuclear Magnetic Resonance (N.M.R) Spectroscopy
Nuclear Magnetic Resonance analyses the behaviour of the nuclei of different atoms under the conditions of the spectrometer. In this way it is very different to Mass Spectrometry, as instead of measuring the time of flight and consequent mass of an isotopic ion such as 2H, N.M.R focuses on the behaviour of the nuclei of atoms. The nuclei of the sample being analysed show magnetic properties when they are placed in a magnetic field. Some of the nuclei line up with the magnetic field and some line up against the magnetic field. Nuclei that line up against the magnetic field have a higher energy than those is aligned with the magnetic field. If the correct frequency of radiation is applied, some nuclei will move up to the next energy level and absorb energy of this frequency. The energy which they absorb corresponds to radio frequency.
The energy/frequency that a nucleus needs to absorb to move up to a higher energy level depends on the strength of the magnetic field that it is in contact with. This magnetic field is affected by other atoms in the molecule as they each have a magnetic field of their own, which can be strengthened or disrupted. Therefore for every type of molecular arrangement, there is a varying magnetic field produced. The atoms have different energy gaps between their energy levels and subsequently absorb different frequencies of radiation. As a result, they give different N.M.R. absorption peaks and it is possible to identify the number of different types of Hydrogen atoms present in a compound, in the context of the different environments they are located in.
The ‘environment’ an atom is situated in plays a key role in the N.M.R spectrum.
An environment is the group that the atom is found in or attached to within a compound, and can cause varying properties such as differences in boiling/melting point, bond length and bond enthalpy. In N.M.R, a proton in a different environment absorbs different amounts of energy, and gives a different chemical shift.
A simple example can be used to explain:
In Phenol, the Hydrogen atom (proton, is in a Phenol environment, as it is attached to a Benzene ring. It will typically have a chemical shift of 4.5-10.0(30).
In Methanol, the H is attached to a CH3 group and is in a carbon environment, and will typically have a chemical shift of 0.8-1.3(30).
As shown, the two protons chemical shift values vary considerably. This is due to the chemical environment they are situated in, and is the key component in identifying compounds in Nuclear Magnetic Resonance Spectroscopy.
As explained above, Protons in different environments absorb energy of different frequencies. NMR spectroscopy measures these differences relative to a standard substance, and the difference is known as Chemical Shift.
The standard substance is usually tetramethylsilane (TMS)(30). This molecule has 12 protons all with identical environments, therefore only produces a single absorption peak. This peak is given a chemical shift value of 0.
Spectra often show a peak at 0 because some TMS is added to the test compound for calibration purposes.
A simplified diagram of NMR apparatus (31):
Interpreting an NMR spectrum
Running along the bottom of an NMR spectrum is a tetramethylsilane (TMS) reference which is used as a standard reference because it gives a sharp signal well away from most of the substances of interest to chemists. The NMR spectrum is differentiated by the extent at which the signal varies from the TMS reference. For example, the NMR spectrum for Ethanal (CH3CHO) (see below) has a peak at a chemical shift of around 10, and then another peak which is three times higher at a chemical shift of around 2. The peak at 10 is from the CHO part of the molecule whereas the peak at 2 is from the CH3 part of the molecule. The reason the CH3 peak is three times larger than the CHO peak is because it has 3 more hydrogen atoms. The integrated trace goes up in steps which are proportional to the areas of absorption signal. This means that the amount of protons being absorbed each time can be found.
Each peak on the NMR spectrum is due to one or more protons in a particular environment.
The relative area under each peak tells you how many protons are in that environment.
There are two peaks on the diagram above, so there are two environments.
The area ratio is 1:3- so there is 1 proton in the environment at 9.5 Chemical shift to every 3 protons in the other environment.
The line on the diagram on the diagram is the integration trace. The height increases are proportional to the areas under the trace and it shows how many protons are in each environment.
As seen from the diagram above, the Hydrogen atoms in Ethanal are all located in different environments, hence the difference in NMR spectrum.
Samples are dissolved
The sample which is being analysed has to be dissolved in a solvent that has no single protons- or they will show up as peaks on the spectrum and have an impact on the results. Denatured solvents are often used-their hydrogen atoms have been replaced by Deuterium. Deuterium is an isotope of Hydrogen that has two nucleons (a proton and a neutron), so there is no overall spin on the nucleus.
My Aspirin sample is being sent for NMR analysis, so an equipment list and method is not necessary. However, detail can be found in the diagram of the NMR apparatus above.
Risk Assessments of Methods
Synthesising Aspirin (Method 1 and 2)
(1=Unlikely, 10=Very likely)
(1=Not Severe, 10= Very Severe)
Ethanoic anhydride is corrosive – could cause blindness or chemical skin burns in direct contact.
Wear gloves, goggles and a lab coat while handling this chemical
Take care whilst pouring, use funnels when appropriate.
Concentrated Sulfuric (VI) acid/Phosphoric Acid is corrosive – may cause blindness or chemical skin burns in direct contact.
Wear gloves, goggles and a lab coat while handling.
Take care while pouring.
Glacial Ethanoic acid is corrosive – may cause blindness or chemical skin burns in direct contact.
Wear gloves, goggles and a lab coat while handling.
Take care while pouring.
Oven is hot – burns may occur when retrieving aspirin samples.
Ask the lab technician to retrieve samples if possible. Take care when putting hands into oven, wear safety gloves.
Ethanoic Anhydride when reacted produces fumes which may cause harm when inhaled, and damage eyes.
Ensure this reaction is carried out in the fume cupboard, and goggles worn to prevent fumes entering eyes.
Dealing with Glass objects. (funnels, etc) which are easily smashed
Ensure equipment such as Hirsch funnels and burettes are clamped into place to prevent them falling and shattering.
Bunsen Burner involves an open flame, which can cause severe burns to the skin.
When using the Bunsen Burner, make sure it is supported by a heat proof mat, and is left on a visible orange flame when possible.
Recrystallisation of impure Aspirin
(1=Non Severe, 10= Very Severe)
Burns from hot water bath/kettle
Ensure care is taken when pouring water from the kettle, use a funnel where appropriate. Make sure hot water bath is not left unsupervised.
Use of Buchner funnel, can tip over and shatter.
Make sure the funnel is clamped to avoid it falling over when the water vacuum is turned on.
Use of water, may cause spillage.
Make sure spillage is kept to a minimum and immediately clear up any water.
Melting Point of Aspirin
(1=Not Severe, 10= Very Severe)
Burns to the skin from Bunsen Burner flame
When using the Bunsen Burner, make sure it is supported by a heat proof mat, and is left on a visible orange flame when possible.
Capillary tubes are sharp, and may cut the skin
Wear gloves whilst handling capillary tubes, and make sure they are disposed of with glass items immediately after use.
Hot apparatus may cause burns if touched
Allow apparatus to cool before touching it. Leave in a safe place to cool, away from flammable liquids or material.
Capillary tubes shatter and fragment when put under pressure
When viewing the tube, ensure goggles are worn to prevent shrapnel entering the eye if the tube snaps.
Iron (III) Chloride Test for Phenol
(1=Non Severe, 10= Very Severe)
Iron (III) Chloride is toxic, corrosive and acidic. May cause harm on contact with the skin.
Ensure gloves are worn when handling Iron (III) Chloride. Also make sure the workbench is clean after use to prevent the chemical from being spread.
Thin Layer Chromatography (T.L.C)
(1=Not Severe, 10= Severe)
Chromatography solvent is a highly flammable solvent, and can easily catch fire.
Do not use the solvent around an open flame, and keep away from hot areas of the lab.
Glassware (beaker, capillary tubes) may cause harm if broken.
Wear goggles and a lab coat while handling glassware.
Prolonged exposure to UV may cause blindness, skin burns and cancer.
Face UV light away from you and others at all times. Ensure goggles are on when using and do not look directly at the light.
If you are the original writer of this essay and no longer wish to have the essay published on the UK Essays website then please click on the link below to request removal: